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Author: Subject: Any other calcium salt from Calcium Sulfate?
hodges
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[*] posted on 2-9-2019 at 15:08
Any other calcium salt from Calcium Sulfate?


I have a bunch of food-grade fine calcium sulfate and was thinking the other day whether it would be possible to convert it into any other calcium salt (chloride, carbonate, etc.). I'm guessing that the answer is no, at least not with any common water-based laboratory reaction. Since calcium sulfate is insoluble, it would not react with anything else.

It would be pretty trivial to go the other way and get calcium sulfate or other calcium salts starting with the chloride (mix with a sulfate and filter out the precipitate), or carbonate (treat with the acid of choice). But going the other way does not seem possible.

Of course I can think of some exotic ways to do it. For example, reducing it at high temperatures to the sulfite or even sulfide in a thermite-type reaction and then reacting the reduced product with the acid of choice. But I could not think of "ordinary" and economic way to get any other calcium salt. Did I miss anything?

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fusso
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[*] posted on 2-9-2019 at 15:49


Yes. Mixing it with Na2CO3 in water will give CaCO3. BaCO3 can be made the same way.
Edit: added link
https://www.sciencemadness.org/whisper/viewthread.php?tid=80...

[Edited on 190903 by fusso]




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happyfooddance
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[*] posted on 2-9-2019 at 19:50


Yup, I do this routinely (metathesis with sodium carbonate) and it is quantitative. But because of the slight solubility of the sulfate it requires some time, heat, and good stirring though.
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[*] posted on 3-9-2019 at 00:23


this is quite something, i always felt like there was a cool use for broken gypsum walls, seeing it as washing powder is basically crude Na2CO3, now were heading for bucket-chemistry



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[*] posted on 3-9-2019 at 06:21


Just out of curiosity what would be the benefit of Na2SO4 or CaCO3, the latter bein available by the truckload for next to nothing and the former, well, not really all that useful or hard to get.
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[*] posted on 3-9-2019 at 09:48


Quote: Originally posted by Antiswat  
this is quite something, i always felt like there was a cool use for broken gypsum walls, seeing it as washing powder is basically crude Na2CO3, now were heading for bucket-chemistry


In the case of a zombie apocalypse, insulation from abandoned buildings is going to be our preferred supplier of gypsum, and hence CaCO3/CaO... This fact had not escaped me :)
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[*] posted on 3-9-2019 at 13:04


You can reduce it with characol to CaS wenn you added Fe2O3 this Reaktion happens at 800C and takes 1/2 to 1H.
When you use 2 parts CaSO4 on one part Fe2O3 and heat it to bright yellow you get SO3.
P.S i know it s not gyps but when you have lots of CaCO3 you can mix it with NaCL. After time you get on the surface Soda and on the Bottom CaCl solution. Thats the way nature makes so Soda sees.
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hodges
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[*] posted on 3-9-2019 at 16:14


Quote: Originally posted by happyfooddance  
Yup, I do this routinely (metathesis with sodium carbonate) and it is quantitative. But because of the slight solubility of the sulfate it requires some time, heat, and good stirring though.


So you are saying that even though CaSO4 has a very low solubility, CaCO3 is even lower? How would one separate the produced CaCO3 from CaSO4? Just use stoichiometric amounts and wait a long time to be sure the reaction is complete, then wash the product?

CaCO3 could then easily be converted to CaCl2 with dilute HCl.

Might be worth trying to see if this works. I should be able to tell if I ended up with CaCl2 (instead of still having CaSO4 mixed with Na2CO3) by testing if the final compound is hygroscopic. Actually, if the product dissolves completely upon addition of HCl that should prove it since CaSO4 is presumably not soluble in HCl.
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[*] posted on 3-9-2019 at 16:36


Quote: Originally posted by hodges  
Quote: Originally posted by happyfooddance  
Yup, I do this routinely (metathesis with sodium carbonate) and it is quantitative. But because of the slight solubility of the sulfate it requires some time, heat, and good stirring though.


So you are saying that even though CaSO4 has a very low solubility, CaCO3 is even lower? How would one separate the produced CaCO3 from CaSO4?


There is quite a big difference in the solubility of the 2, CaSO4 deposits I can usually remove with hot water, CaCO3...not so much. I just use a small (10% maybe, I actually just use molar weights but my calcium sulfate has maybe 10% moisture) excess of sodium carbonate, and there is no sulfate left after 2 hrs of heat and stirring (3-4 mol, bucket scale).

The sodium sulfate/carbonate solution is discarded and the carbonate is washed with water until rinse water is only very slightly basic. This is most easily accomplished by letting the carbonate settle for a few days and then decanting.

I think there is a rough description in the Golden Book of Chemistry (in the forum library).

Edit: Actually, the part in the GBC I was thinking of was making NaOH from CaO and sodium carbonate.

But it works the same.

[Edited on 9-4-2019 by happyfooddance]
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[*] posted on 4-9-2019 at 01:18


Well, just for the sake of completeness, I think it is possible to boil CaSO4 with sulfur and water to have some kind of disproportionation, or try to melt sulfur with the sulfate. There are a little chances you will get some calcium thiosulfate this way which is very soluble in water.

Another option could be conversion to calcium bisulfate (+H2SO4) which is much more soluble (according to Treadwell).

Also from Treadwell:

"Calcium sulfate is also soluble in concentrated ammonium sulfate, owing the formation of a complex ion:

CaSO4 + (NH4)2SO4 = (NH4)2[Ca(SO4)2]"



[Edited on 4-9-2019 by teodor]

[Edited on 4-9-2019 by teodor]
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[*] posted on 4-9-2019 at 07:21


I remembered posting about this before, and there are useful references in this thread:

http://www.sciencemadness.org/talk/viewthread.php?tid=80545#...
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[*] posted on 4-9-2019 at 15:37


Quote: Originally posted by happyfooddance  
I remembered posting about this before, and there are useful references in this thread:

http://www.sciencemadness.org/talk/viewthread.php?tid=80545#...


Thanks, happyfooddance! I think I will get some washing soda (or perhaps make sodium carbonate from the bicarbonate), make a solution with it in excess in say a 2 liter pop bottle, then add some calcium sulfate and let it sit for a week or so (shaking from time to time). Then I will filter and wash the resulting insoluble product and see if it dissolves fully (with production of CO2) upon adding dilute HCl drop-wise until CO2 production stops. I can then heat it to dry, and see if it becomes wet again upon standing.


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[*] posted on 4-9-2019 at 22:17


Be careful because moist CaCl2 decomposes and evolves HCl, I've ruined a microwave or two this way...

You can minimize the decomposition by adding conc. HCl or drying in a stream of hot HCl gas, but you will still have to deal with the HCl.

Edit: to the credit of fusso, he linked to that thread before I even saw this one:
Quote: Originally posted by fusso  
Yes. Mixing it with Na2CO3 in water will give CaCO3. BaCO3 can be made the same way.
Edit: added link
https://www.sciencemadness.org/whisper/viewthread.php?tid=80...

[Edited on 190903 by fusso]


[Edited on 9-5-2019 by happyfooddance]
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[*] posted on 5-9-2019 at 09:06


This discussion is relevant to some fun I had with some kids, evaporating sea water.

CaSO4 has low solubility but it is not zero, and it is much higher than the solubility of CaCO3. We started with about 60kg of sea water from the Pacific Ocean, first filtered it through coffee filters to remove sand etc. You can evaporate about 1/2 of the water before anything crystallizes out, but at that point you start to get CaSO4. Before that happens we learned it was good advice to filter the water again through activated charcoal, to remove organic matter.

Between about 50% evaporation and 90% you get nice crystals of CaSO4, they are very pretty, kind of milky in color. You also get some CaCO3, which is in solution because of dissolved CO2 (that is, you have calcium bicarbonate, which precipitates as CaCO3 as the CO2 is lost).

At about 90% evaporation you start to get NaCl crystals, but CaSO4 continues to crystallize as well. The latter can be separated mechanically with tweezers, as the crystals are obviously quite different from NaCl crystals.

We separated our CaSO4 crystals from the CaCO3 by washing in a little HCl. I always assumed that by adding Na2CO3 to the seawater in the first place you could convert all the CaSO4 into CaCO3 which would then precipitate, thereby removing all calcium , but I never tried this. There is also a small amount of strontium. SrSO4 is much less soluble than CaSO4 so there's a question of why it's there at all. The answer seems to be that the solubility of SrSO4 is affected by the presence of other ions. I never did figure out a good way of separating the strontium, but there isn't that much of it anyway.




Any other SF Bay chemists?
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[*] posted on 7-9-2019 at 13:42


I did try this, and verify that it worked.

To 500 ml of distilled water, I slowly added CaSO4 until it failed to dissolve. That happened at the 1 gram point - would estimate less than 0.1g did not dissolve.

I heated several ounces of NaHCO3 in a Pyrex cup in the oven at 400 degrees (F) for just over an hour. Based on weighing before and after, I concluded that only about 85% of the NaHCO3 was converted to Na2CO3 in this amount of time. I put it back in the oven, this time at 425 F, for another hour, and calculations showed that somewhere around 92% had been converted.

I dissolved several grams the Na2CO3 in a few ounces of distilled water, and mixed with the CaSO4 solution. CaCO3 immediately precipitated. I shook the solution and left it overnight. Next day I carefully poured off most of the water, then again filled to 500ml (with tap water), and shook. After that settled, I again poured off most of the water.

I added some universal indicator to the remaining water and product and shook. Based on the color, pH was around 9. I than proceeded to add dilute HCl (slowly so as not to add too much), noting that bubbles formed as expected, until the pH reached about 6. I then poured this into a saucer and dried in an oven at 425 F (took about 1/2 hour).

Unfortunately, I was unable to weight the product. As soon as I took it out of the oven, it began to become moist, and was shiny with water within a few minutes. But this deliquescence does show that the product was likely mostly CaCl2.

I have two tiny plastic saucers into which I weighed out 1.00g of the Na2CO3 I made (presumably anhydrous). These are sitting in the room, normal indoor temperature and humidity. The first I left dry, while the second I dissolved in a minimum amount of water. Presumably, they should both converge to the same weight, representing a hydrate of somewhere between 1 and 10 H2O. The dry one so far (one day) has gained under 0.05 grams. The wet one still has not started to crystallize.



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[*] posted on 9-9-2019 at 01:27


how to extract CaCO3? i experimented with dissolving CaCO3 using H2CO3, carbonated water. it works. i was stunned, it converts the CaCO3 into CaHCO3, this is soluble, you then decompose this and out of seemingly nowhere you get CaCO3 ppt'ing
chunks of dry ice would be quite neat for extracting CaCO3, i believe solubility is on solubility table even, CaHCO3 that is.
i believe it was decomposing due to high temperature as i was boiling it down, wonder what CaHCO3 crystals look like.

interestingly, CaSO4 is soluble in H2SO4, but it appears that it has to be somewhere mid concentration of sulfuric acid, i really dont like playing around with sulfuric acid, i happened to accidentally dissolve calcium sulfate in sulfuric acid i had in a plastic bottle, waste acids from whatever reactions i would use to clean limescale in toilet, spikey crystals. the CaSO4 crystals formed as i would dilute the acid mixture down further and further so there must be a very particular concentration range in which CaSO4 is soluble in H2SO4




~25 drops = 1mL @dH2O viscocity - STP
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https://en.wikipedia.org/wiki/Solubility_table
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[*] posted on 11-9-2019 at 13:08


For completeness, though a tangent from the original preparation, here is what I found for equilibrium water content of the Na2CO3, at a temperature around 23C, and relative humidity near 55%.

Na2CO3 has a molecular weight of 106, with water having a molecular weight of 18. Therefore, each water molecule will add 18 / 106 = 17% to the mass. I started with two 1.00g (presumably anhydrous) Na2CO3 samples. One I added enough water to fully dissolve. The mass is now stable for both samples. The dry sample weighs 1.11g, and the sample with evaporated water weighs 1.19g. So I conclude there is one molecule of water for every sodium carbonate molecule.

If commercial washing soda really does have .10H2O, well over half the mass can be expected to effervesce when leaving the open box.
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[*] posted on 2-12-2019 at 04:32


i have experienced washing powder clumping up a bit after longer time, and since its granules its quite a bit more hygroscopic than if same effect was seen with powder, at least the brand i was using likely wasnt 10H2O, it would from an economic perspective make sense for them to add more water to sell their product for same price, but with less actual product



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